A Level Notes Gp5

7/27/2019 A Level Notes Gp5

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Group 5/15 Introduction

down group 5/15 ===>

property\Zsymbol, name 7N Nitrogen 15P Phosphorus 33As Arsenic 51Sb Antimony 83Bi Bismuth

Period 2 3 4 5 6

Appearance (RTP) colourless gaswhite/red solid

allotropes

grey solid (also

yellow/black

allotropes)

grey metalloid

solid(also yellow

allotrope)

silver-white brittle metal

melting pt./oC -210 44 sublimes 631 272

boiling pt./oC -1.96 280 616? 1635 1560

density/gcm-3


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principal oxidation states -3 to +5 -3, +3, +5 +3, +5 +3, +5 +3, (+5)

property\Zsymbol, name 7N Nitrogen 15P Phosphorus 33As Arsenic 51Sb Antimony 83Bi Bismuth

Generally speaking down a p block group the element becomes more metallic in chemical character.

Nitrogen and phosphorus are non-metals, arsenic and antimony are semi-metals, bismuth is a true metal.

NITROGEN - brief summary of a few points

The structure of the element:

o Non-metal existing as diatomic molecule N2, with a triple covalent bond, N N.

Physical properties:

o Colourless gas; mpt -210oC; bpt -196

oC; poor conductor of heat/electricity.

Group, electron configuration (and oxidation states):

o Gp5; e.c. 2.5 or1s22s

22p

3; Variety of oxidation states from -3 to +5 e.g.

o NH3 (-3), N2O (+1), NO (+2), NCl3 (+3), NO2 (+4) and N2O5 and HNO3 (+5).

Reaction of nitrogen with oxygen:

o At high temperatures e.g. in car engines, nitrogen(II) oxide (nitrogen monoxide) is formed.

N2(g) + O2(g) ==> 2NO(g)

o and the nitrogen(II) oxide rapidly reacts in air to form nitrogen(IV) oxide (nitrogen dioxide).

2NO(g) + O2(g)==> 2NO2(g)
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o The theoretical highest oxide is N2O5 nitrogen(V) oxide (nitrogen pentoxide) and does exist.

Reaction of nitrogen oxides with water:

o Nitrogen(IV) oxide dissolves to form an acidic solution of weak nitrous acid and strong nitric acid.

2NO2(g) + H2O(l)==> HNO2(aq) + HNO3(aq)

or2NO2(g) + 2H2O(l) ==> HNO2(aq) + H3O+

(aq) + NO3-(aq)

o NO and N2O are neutral oxides but nitrogen(V) oxide is strongly acidic and dissolves to form nitric acid.

N2O5(s) + H2O(l)==> 2HNO3(aq)

orN2O5(s) + 3H2O(l)==> 2H3O+

(aq) + 2NO3-(aq)

Reaction of nitrogen oxides with acids:

o None, only acidic (N2O3 (very unstable), NO2 and N2O5) or neutral (N2O and NO), in nature.

Reaction of nitrogen oxides with bases/alkalis:

o Nitrogen(IV) oxide or nitrogen dioxide forms sodium nitrite and sodium nitrate with sodium hydroxide solution.

o 2NO2(g) + 2NaOH(aq) ==> NaNO2(aq) + NaNO3(aq) + H2O(l)

o ionic equation: 2NO2(g) + 2OH-(aq) ==> NO2

-(aq) + NO3

-(aq) + H2O(l)

o As well as being a neutralisation reaction, it is also a redox reaction, the oxidation states of oxygen (-2) and hydrogen (+1) do not

change BUT the oxidation state of nitrogen changes from two at (+4) to one at (+3) and one at (+5). The simultaneous change of

an element into an lower and upper oxidation sate is sometimes called disproportionation.

Reaction of nitrogen with chlorine:

o None, but the unstable yellow oily liquid chloride can be made indirectly.

Reaction of chloride with water:

o Slowly hydrolyses to form weak nitrous acid and strong hydrochloric acid.

NCl3(l) + 2H2O(l) ==> HNO2(aq) + 3HCl(aq)

orNCl3(l) + 2H2O(l) ==> HNO2(aq) + 3H+

(aq) + 3Cl-(aq)

orNCl3(l) + 5H2O(l) ==> HNO2(aq) + 3H3O+

(aq) + 3Cl-(aq)

Reaction of nitrogen with water:


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o Slightly soluble but NO reaction.

Other comments:

o An essential element for plants, hence need for nitrogen compounds in compost and artificial fertilisers (NPK

bags!).

PHOSPHORUS - brief summary of a few points

The structure phosphorus:

o Two solid allotropes (red and white) consisting ofP4 molecules, also a polymer form.

Physical properties of phosphorus:

o Colourless gas; mpt 44oC; bpt 280

oC; poor conductor of heat/electricity.

Group, electron configuration (and oxidation states):

o Gp5; e.c. 2,8,5 or1s22s

22p

63s

23p

3; Variety of oxidation states from -3 to +5 e.g.

o PH3 (-3), P4O6 (+3), P4O10, PCl5 and H3PO4 (+5).

Reaction of phosphorus with oxygen:o With limited air/oxygen, on heating the phosphorus, the covalent white solid phosphorus(III) oxide is formed.

P4(s) + 3O2(g) ==> P4O6(s)

o With excess air/oxygen, on heating the phosphorus, the covalent white solid phosphorus(V) oxide is formed.

P4(s) + 5O2(g) ==> P4O10(s)

Reaction of phosphorus oxides with water: Both oxides dissolve in water to form acidic solutions.
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o Phosphoric(III) oxide forms phosphoric(III) acid.

P4O6(s) + 6H2O(l) ==> 4H3PO3(aq)

o Phosphoric(III) oxide forms phosphoric(V) acid.

P4O10(s) + 6H2O(l) ==> 4H3PO4(aq)

Reaction of phosphorus with chlorine:

o With limited chlorine, on heating the phosphorus, the covalent liquid phosphorus(III) chloride is formed.

P4(s) + 3Cl2(g) ==> 4PCl3(l)

o With excess chlorine, on heating the phosphorus, the ionic*solid phosphorus(V) chloride is formed.

P4(s) + 5Cl2(g) ==> 4PCl5(s)

*PCl5 is a bit unusual for an 'expected covalent' liquid chloride.

It is an ionic solid with the structure [PCl4]+[PCl6]

-

Hence its melting point is much greater than the liquid phosphorus(III) chloride, where the molecules are

only held together by the inter-molecular forces.

However, gaseous phosphorus(V) chloride consists ofPCl5 covalent molecules.

Reaction of phosphorus oxides with acids:

o None, only acidic in nature.

Reaction of phosphorus oxides with bases/alkalis:

o Both oxides dissolve in alkalis to form a whole series of phosphate(III) and phosphate(V) salts.

o with excess strong bases like sodium hydroxide, the simplified equations are:

P4O6(s) + 12NaOH(aq) ==> 4Na3PO3(aq) + 6H2O(l) sodium phosphate(III) formed from phosphorus(III) oxide

ionic equation: P4O6(s) + 12OH-(aq) ==> 4PO3

3-(aq) + 6H2O(l)

P4O10(s) + 12NaOH(aq) ==> 4Na3PO4(aq) + 6H2O(l) sodium phosphate(V) formed from phosphorus(V) oxide

ionic equation: P4O10(s) + 12OH-(aq) ==> 4PO4

3-(aq) + 6H2O(l)

If the empirical formulae P2O3 and P2O5 are used, just halve all the balancing numbers.

Other than using excess sodium hydroxide, other salts can be formed.


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e.g. P4O10(s) + 4NaOH(aq) + 2H2O(l) ==> 4NaH2PO4(aq) sodium dihydrogen phosphate(V)

or P4O10(s) + 8NaOH(aq) ==> 4Na2HPO4(aq) + 2H2O(l) disodium hydrogen phosphate(V)

o -

Reaction of phosphorus chlorides with water:

o Phosphorus(III) chloride hydrolyses rapidly and exothermically to form phosphoric(III) acid.

PCl3(l) + 3H2O(l) ==> H3PO3(aq) + 3HCl(aq)

o Phosphorus(V) chloride initially hydrolyses to form phosphorus oxychloride? and hydrochloric acid.

PCl5(s) + H2O(l) ==> POCl3(aq) + 2HCl(aq)

Then on boiling the aqueous solution, phosphoric(V) acid is formed and more hydrochloric acid.

POCl3(aq) + 3H2O(l) ==> H3PO4(aq) + 3HCl(aq)

overall: PCl5(s) + 4H2O(l) ==> H3PO4(aq) + 5HCl(aq)

Reaction of element with water:

o None.

The shapes of some molecules and ions of nitrogen and phosphorus


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electrons: three bond pairs and one lone pair, PYRAMIDAL or TRIGONAL PYRAMID shape: e.g. ammonia NH3 with bond angle of

approximately 109o. Note: the exact H-N-H angle is 107

odue to the extra repulsion of one lone pair (for H-X-H angles: NH3 > H2O and BF3Boron trifluoride (3 bonding pairs, 6 outer electrons) acts as a lone pair acceptor (Lewis acid) and ammonia (3 bond pairs)

and lone pair which enables it to act as a Lewis base - a an electron pair donor. It donates the lone pair to the 4th 'vacant' boron orbital to form

a sort of'adduct' compound. Its shape is essentially the same as ethane, a sort of double tetrahedral with H-N-H, N-B-F and

F-B-F bond angles of ~109o.
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Nitrogen(IV) oxide, NO2 (nitrogen dioxide) is bent shaped (angular), O-N-O

bond angle ~120obecause of two bonding groups of bonding electrons

and a single lone electron in the same plane as the bonding pairs of

electrons.

The nitrate(III) ion, NO2-(nitrite ion) is bent shaped (angular), O-N-O bond

angle ~120o

due to two groups of bonding electrons and one lone pair of

electrons.

The nitrate(V) ion, NO3-(nitrate ion) is trigonal planar, O-N-O bond angle

120o

due to three bonding groups of electrons and no lone pairs of

electrons.

The nitronium ion, NO2+, is linear, O-N-O bond angle of 180

obecause

there are two groups of bonding electrons and no lone pairs of

electrons (you easily see this from the NO2 neutral molecule diagram below).

The shapes are deduced below using dot and cross diagrams and VSEPR

theory and illustrated in the valence bond dot and cross diagrams below.

Ammonia can act as an electron pair donor ligand in transition metal ion complexes

e.g. tetraamminedichlorochromium(III) complex ion - cis/trans isomers (Z/E isomers)


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With excess aqueous ammonia a pale blue hexa-ammine complex is formed with hexaaquanickel(II) ions

[Ni(H2O)6]2+

(aq) + 6NH3(aq) [Ni(NH3)6]2+

(aq) + 6H2O(l)

and also excess aqueous ammonia a pale blue hexa-ammine complex is formed with aqueous copper(II) ions

[Cu(H2O)6]2+

(aq) + 4NH3(aq) [Cu(NH3)4(H2O)2]2+

(aq) + 4H2O(l)

or[Cu(H2O)4]2+

(aq) + 4NH3(aq) [Cu(NH3)4]2+

(aq) + 4H2O(l)

The Synthesis of ammonia - The Haber Process
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Ammonia gas is synthesisedin the chemical industry by reacting nitrogen gas with hydrogen gas in what is known as the Haber-

Bosch Process, named after two highly inventive and subsequently famous chemists.

The nitrogen is obtained from liquified air(80% N2). Air is cooled and compressed under high pressure to form liquid air (liquefaction).The liquid air is fractionally distilled at low temperature to separate oxygen (used in welding, hospitals etc.), nitrogen (for making

ammonia), Noble Gases e.g. argon for light bulbs, helium for balloons).

The hydrogen is made by reacting methane (natural gas) and wateror from cracking hydrocarbons (both reactions are done at high

temperature with a catalyst).

o CH4 + H2O ==> 3H2 + CO

o e.g. C8H18==> C8H16 + H2

The synthesis equation for this reversible reaction is ...o N2(g) + 3H2(g) 2NH3(g)

.. which means an equilibrium will form, so there is no chance of 100% yield even if you use, as you actually do, the theoretical reactant

ratio of nitrogen : hydrogen of 1 : 3

In forming ammonia 92kJ of heat energy is given out (i.e. exothermic, 46kJ of heat released per mole of ammonia formed).

Four moles of 'reactant' gas form two moles of 'product' gas, so there is net decrease in gas molecules on forming ammonia.

So applying the Le Chatelier equilibrium rules, the formation of ammonia is favoured by ...

o (a) Using high pressure because you are going from 4 to 2 gas molecules (the high pressure also speeds up the reaction

because it effectively increases the concentration of the gas molecules), but higher pressure means more dangerous and more

costly engineering.

o (b) Carrying out the reaction at a low temperature, because it is an exothermic reaction favoured by low temperature, but this

may produce too slow a rate of reaction,

o So, the idea is to use a set ofoptimum conditions to get the most efficient yield of ammonia and this involves getting a low %

yield (e.g. 8% conversion) but fast. Described below are the conditions to give the most economic production of ammonia.

o these arguments make the point that the yield* of an equilibrium reaction depends on the conditions used.


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* The word 'yield' means how much product you get compared to the theoretical maximum possible if the

reaction goes 100%.

For more on chemical economics seeExtra Industrial Chemistrypage.

In industry pressures of200 - 300 times normal atmospheric pressure are used in line with the theory.

Theoretically a low temperature would give a high yield of ammonia BUT ...

o Nitrogen is very stable molecule and not very reactive i.e. chemically inert, so the rate of reaction is too slow at low

temperatures.

o To speed up the reaction an iron catalyst is used as well as a higher temperature (e.g. 400-450oC).

o The higher temperature is an economic compromise, i.e. it is more economic to get a low yield fast, than a high yield slowly!

o Note: a catalyst does NOT affect the yield of a reaction, i.e. the equilibrium position BUT you do get there faster!

At the end of the process, when the gases emerge from the iron catalyst reaction chamber, the gas mixture is cooled under high

pressure, when only the ammonia liquefies and is so can be removed and stored in cylinders.

Any unreacted nitrogen or hydrogen (NOT liquified), is recycled back through the reactor chamber, nothing is wasted! [nitrogen (-

196oC) and hydrogen (-252

oC) have much lower boiling points than ammonia (-33

oC). Boiling points increase with pressure, but these

normal atmospheric pressure values offer a fair comparison].

To sum up: A low % yield of ammonia is produced quickly at moderate temperatures and pressure, and is more economic than

getting a higher % equilibrium yield of ammonia at a more costly high pressure and a slower lower temperature reaction.

AND there are some more general notes onChemical Economicson the Industrial Chemistry page.
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The Uses of Ammonia and derived compoundsAmmonia is used to manufacture nitric acid

Ammonia is oxidised with oxygen from air using a hot platinum catalyst to form nitrogen monoxide and water.

4NH3(g) + 5O2(g)==> 4NO(g) + 6H2O(g)

The gas is cooled and reacted with more oxygen to form nitrogen dioxide.

2NO(g) + O2(g) ==>2NO2(g)

This is reacted with more oxygen and water to form nitric acid.

4NO2(g)+ O2(g) + 2H2O(l)==> 4HNO3(aq)

Nitric acid is used to make nitro-aromatic compounds from which dyes are made.

It is also used in the manufacture of artificial nitrogenous fertilisers(like ammonium nitrate, see below).

Ammonia is used to manufacture 'artificial' nitrogenous fertilisers

Ammonia is a pungent smelling alkaline gasthat is very soluble in water.

The gas or solution turns litmus or universal indicator bluebecause it is a soluble weak base or weak alkali and is neutralised by

acids to form salts.

Ammonium salts are used as 'artificial' or 'synthesised' fertilisers i.e. nitrogenous fertilisers 'man-made' in a chemical works, and

used as an alternative to natural manure or compost etc.

The fertiliser salts are made by neutralising ammonia solution with the appropriate acid.

The resulting solution is heated, evaporating the water to crystallise the salt e.g.

o ammonia + sulphuric acid ==> ammonium sulphate

o 2NH3(aq) + H2SO4(aq)==> (NH4)2SO4(aq)


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o AND

o ammonia + nitric acid ==> ammonium nitrate

o NH3(aq) + HNO3(aq)==> NH4NO3(aq)

The salt Ammonium chloride is used in zinc-carbon dry cell batteries. The slightly acid paste, made from the salt, slowly reacts with the

zinc to provide the electrical energy from the chemical reaction.

o ammonia + hydrochloric acid ==> ammonium chloride

o NH3(aq) + HCl(aq)==> NH4Cl(aq)

orNH4OH(aq) + HCl(aq)==> NH4Cl(aq) + H2O(l)

If ammonium salts are mixed with sodium hydroxide solution, free ammonia is formed(detected by smell and damp red litmus

turning blue).

o

e.g. ammonium chloride + sodium hydroxide ==> sodium chloride + water + ammonia

o NH4Cl + NaOH ==> NaCl + H2O + NH3

Ammonium sulphate or nitrate salts are widely used as 'artificial orsynthetic fertilisers (preparation reactions

above). There are several advantages to using artificial fertilisers in the absence of sufficient manure-silage etc. e.g. relatively cheap

mass production, easily used to make poor soils fertile or quickly enrich multi-cropped fields.

Artificial fertilisers are important to agriculture and used on fields to increase crop yields but they should be applied in a

balanced manner.

o Fertilisers usually contain compounds ofthree essential elements for healthy and productive plant growth to increase crop

yield. They replace nutrient minerals used by a previous crop or enriches poor soil and more nitrogen gets converted into plantprotein.

Nitrogen (N) e.g. from ammonium or nitrate salts like ammonium sulphate, ammonium sulphate or ammonium phosphate

or urea (e.g. look for the N in the formula of ammonium salts)

Phosphorus (P) e.g. from potassium phosphate or ammonium phosphate

Potassium (K) e.g. from potassium phosphate, potassium sulphate.


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The fertiliser is marked with an 'NPK' value, i.e. the nitrogen : phosphorus : potassium ratio

o Fertilisers must be soluble in water to be taken in by plant roots.

Problems with using 'artificial' fertilisers Overuseof ammonia fertilisers on fields can cause major environmental problems as well as being uneconomic.

Ammonium salts are water solubleand get washed into the groundwater, rivers and streams by rain contaminating them with

ammonium ions and nitrate ions.

This contamination causes several problems.

Excess fertilisers in streams and rivers cause eutrophication.

o

Overuse of fertilisers results in appreciable amounts of them dissolving in rain water.

o This increases levels of nitrate or phosphate in rivers and lakes.

o This causes 'algal bloom' i.e. too much rapid growth of water plants on the surface where the sunlight is the strongest.

o This prevents light from reaching plants lower in the water.

o These lower plants decay and the active aerobic bacteria use up any dissolved oxygen.

o This means any microorganisms or higherlife forms relying on oxygen cannot respire.

o All the eco-cycles are affected and fish and other respiring aquatic animals die.

o The river or stream becomes 'dead' below the surface as all the food webs are disrupted.

Nitrates are potentially carcinogenic(cancer or tumour forming).o The presence in drinking water is a health hazard.

o Rivers and lakes can be used as initial sources for domestic water supply.

o You cannot easily remove the nitrate from the water, it costs too much!

o So levels of nitrate are carefully monitored in our water supply.


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Cost - The hydrogen for the Haber Process for manufacturing ammonia is usually obtained from hydrocarbon sources e.g. methane gas.

Therefore, as oil becomes more scarce, the cost of producing 'artificial' fertilisers will increase.

The Nitrogen Cycle for the gaseous element N2(g)

Nitrogen is an extremely important element for all plant or animal life! It is found in important molecules such as amino acids, which

are combined to form proteins. Protein is used everywhere in living organisms from muscle structure in animals to enzymes in

plants/animals.

Nitrogen from the atmosphere:

o Nitrifying bacteria, e.g. in the root nodules of certain plants like peas/beans (the legumes), can directly convert atmospheric

nitrogen into nitrogen compounds in plants e.g. nitrogen => ammonia => nitrates which plants can absorb.

However, most plants can't do this conversion from nitrogen => ammonia, though they can all absorb nitrates, so the

'conversion' or 'fixing' ability might be introduced into other plant species by genetic engineering.

o The nitrogen from air is converted into ammonia in the chemical industry, and from this artificial fertilisers are manufactured to

add to nutrient deficient soils. However, some of the fertiliser is washed out of the soil and can cause pollution.

o The energy oflightning causes nitrogen and oxygen to combine and form nitrogen oxides which dissolve in rain that falls on the

soil adding to its nitrogen content.

1. N2(g) + O2(g)==> 2NO(g), then


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2. then 2NO(g) + O2(g)==> 2NO2(g)

3. NO2(g) + water==> nitrates(aq) in rain/soil

4. Incidentally, reactions 1. and 2. can also happen in a car engine, and NO2 is acidic and adds to the polluting acidity ofrain as well as providing nutrients for plants!

Nitrogen recycling apart from the atmosphere:

o Nitrogen compounds, e.g. protein formed in plants or animals, are consumed by animals higher up the food chain and then

bacterial and fungal decomposersbreak down animal waste and dead plants/animals to release nitrogen nutrient compounds into

the soil (e.g. in manure/compost) which can then be re-taken up by plants.

Nitrogen returned to the atmosphere:

o However, denitrifying bacteria will break down proteins completely and release nitrogen gas into the

atmosphere.

Some nitrogen oxides chemistry

The equilibria between oxygen O2, nitrogen (II) oxide NO, nitrogen(IV) oxide NO2 and its dimerN2O4.

NO2 can be made from the irreversible thermal decomposition of lead(II) nitrate in a pyrex boiling tube connected to a 100 cm3

gas

syringe in a fume cupboard.

o lead(II) nitrate ==> lead(II) oxide + nitrogen(IV) oxide + oxygen


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o 2Pb(NO3)2(s) ==> 2PbO(s) + 4NO2(g) + O2(g)

(a) 2NO2(g, brown) 2NO(g, colourless) + O2(g, colourless) (H = +113 kJ mol-1

)

o (a) The temperature effect can be observed by strongly heating the gases in the pyrex tube above 400oC.

(b) 2NO2(g, brown) N2O4(g, colourless) (H = -58 kJ mol-1

)

o (b) The temperature effect can be observed by cooling and warming below 100oC.

o (b) The pressure effect can be observed by sealing the cool gases in the gas syringe and compressing and decompressing it.

Temperature and energy change (H)

o (a) Increases in temperature favours the endothermic decomposition of NO2 to NO and O2, so at high temperatures the brown

colour fades.

o (b) Decrease in temperature favours the exothermic formation of the dimer N2O4 from NO2, so the brown colour fades on cooling

the gas mixture.

Gas pressure change (V)

o (a) Increase in pressure favours the LHS, more NO2, because 2 mol gas 1 mol gas, so the mixture would get lighter in colour.

(b) This can be demonstrated by compressing/decompressing the gas mixture in the syringe to see the brown colour

intensity increase/decrease.

In fact you can even see the dynamic equilibrium 'kinetics' in operation here. There is a time lag of about 1-2 seconds

before the new equilibrium position is established as the 'imposed' colour intensity change becomes constant.

Concentration change

o (a) Theoretically an increase in O2 would lead to decrease in NO and increase in NO2, so the mixture would get darker.

o (b) Increase in NO2 would increase N2O4, but overall the colour would still be darker because not all of the 'extra' NO2 can be

converted to maintain the equilibrium.

The formation of nitrogen(II) oxide at high temperature e.g. in a car engine


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o N2(g) + O2(g) 2NO(g)(H = +181 kJ mol-1

)

o Temperature and energy change (H)

Increase in temperature favours the endothermic formation of NO.

This reaction does not happen at room temperature but is formed at the high temperatures in car engines.

Unfortunately when released through the car exhaust, it cools to normal temperatures when NO irreversibly reacts with

oxygen in air to form nitrogen(IV) oxide, NO2,

2NO(g) + O2(g) ==> 2NO2(g)

Nitrogen dioxide is acidic, a lung irritant and a reactive free radical molecule involved in the chemistry of photochemical

smog not good!

o Its concentration in car exhaust gases can be reduced, along with that of carbon monoxide, by using a catalytic converter.

o Using platinum, and other transition metal, based catalysts, the following reaction can be made to take place producing harmless

nitrogen and carbon dioxide.

2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g)

Some acid-base chemistry of ammonia

Definition and examples of WEAK BASES

o A weak base is only weakly or partially ionised in watere.g.
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o A good example is ammonia solution, which is only about 2% ionised :

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

Ammonia is the base and the ammonium ion NH4+

is its conjugate acid.

Water is the acid and the hydroxide ion is its conjugate base.

This equilibrium is sometimes referred to as a base hydrolysis.

The low % of ionisation gives a less alkaline solution of lower pH than for strong soluble bases (alkalis), but pH is

still > 7.

The concentration of water is considered constant and to solve simple problems, the base ionisation equilibrium

expression is written as:

Kb =

[NH4+

(aq)] [OH-(aq)]

-----------------------------

[NH3(aq)]

Kb is the base ionisation/dissociation constant (mol dm-3

) for any base.

Salts of weak bases and strong acids give acidic solutions.

o e.g. ammonium chloride. The chloride ion is such a weak base that there is no acid-base reaction with water, but the ammonium

ion is an effective proton donor. As a general rule, the conjugate acid of a weak base is quite strong. The result here is that

ammonium salt solutions have a pH of 3-4.

o NH4+

(aq) + H2O(l) NH3(aq) + H3O+

(aq)

o In zinc-carbon batteries an acidic ammonium chloride paste dissolves the zinc in the cell reaction, though an oxidising agent must

be added (MnO2) to oxidise the hydrogen formed into water, or batteries might regularly explode!


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o If you place a piece of magnesium ribbon or a zinc granule in ammonium chloride or ammonium sulphate solution you will see

fizzing as hydrogen gas is formed.

2H3O+

(aq) + M(s)==> M2+

(aq) + H2O(l) + H2(g)

M = zinc or magnesium

Buffering action ammonium salts

o A mixture of a weak base and the salt of the weak base with a strong acid.

o e.g. ammonia NH3 and ammonium chloride NH4+Cl

-

o NH4+

and NH3 constitute a conjugate acid-base pair.

o In solution most of the ammonia is NOT ionised (and even suppressed by the ammonium ions from the salt).

It is the weak base that 'removes' most of any added hydrogen ions.

NH3(aq) + H+

(aq) NH4+

(aq)

o The salt is fully ionised in solution giving relatively high concentrations of the ammonium ion.

It is the ammonium ion that removes most of any added hydroxide ions.

NH4+

(aq) + OH-(aq) NH3(aq) + H2O(l)

The thermal decomposition of ammonium chloride
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On heating strongly above 340oC, the white solid ammonium chloride, thermally decomposes into a mixture of two colourless gases

ammonia and hydrogen chloride.

On cooling the reaction is reversed and solid ammonium chloride reforms.

o This is an example ofsublimation but here it involves both physical and chemical changes.

o When a substance sublimes it changes directly from a solid into a gas without melting and on cooling reforms the solid without

condensing to form a liquid.

o Ammonium chloride + heat ammonia + hydrogen chloride

o NH4Cl(s) NH3(g) + HCl(g)o so the thermal decomposition of ammonium chloride is the forward reaction, and the formation of ammonium chloride is the

backward reaction.

Note:

o Reversing the reaction conditions reverses the direction of chemical change, typical of a reversible reaction.

o Thermal decomposition means using 'heat' to 'break down' a molecule into smaller ones.

o The decomposition is endothermic (heat absorbed or heat taken in) and the formation of ammonium chloride is

exothermic (heat released or heat given out).

o This means if the direction of chemical change is reversed, the energy change is also reversed.

o Ammonium fluoride (>?oC), ammonium bromide (>450

oC) and ammonium iodide (>550

oC), with a similar formula, all

sublime in a similar physical-chemical way when heated, so the equations will be similar i.e. just swap F, Br or I for the Cl.

Similarly, ammonium sulphate also sublimes when heated above 235oC and thermally decomposes into ammonia gas

and sulphuric acid vapour.

(NH4)2SO4(s) NH3(g) + H2SO4(g)

Ammonium nitrate does not undergo a reversible sublimation reaction, it melts and then decomposes into

nitrogen(I) oxide gas (dinitrogen oxide) and water vapour.


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NH4NO3(s) N2O(g) + 2H2O(g)

This is very different reaction, in fact its an irreversible redox reaction. The nitrate ion, NO3-, or any nitric acid formed,

HNO3, act as an oxidising agent and oxidise the ammonium ion. If the products are cooled, ammonium nitrate is NOT

reformed.

o For more on sublimation, see theStates of Matterwebpage.

REDOX analysis of selected reactions

The oxidation of ammonia with molecular oxygen

The concept of oxidation state can now be fully applied to reactions which do not involve ions e.g.

The oxidation of ammonia via a Pt catalyst at high temperature which is part of the chemistry of nitric acid manufacture.

4NH3(g) + 5O2(g) ==> 4NO(g) + 6H2O(g) The oxidation number analysis is:

o 4N at (-3) each in NH3 and 10O all at (0) in O2 change to ...

o 4N at (+2) each in NH3, 4O at (-2) each and 6 O at (-2) each in H2O.

o H is +1 throughout i.e. does not undergo an ox. state change.

o Oxygen is reduced from ox. state (0) to (-2).
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o Nitrogen is oxidised from ox. state (-3) to (+2).

o The total increase in ox. state change of 4 x ( -3 to +2) for nitrogen is balanced by the total decrease in ox. state change of 10 x (0

to -2) for oxygen i.e. 20 e-or ox. state units change in each case.

o Oxygen is the oxidising agent (gain/accept e-s, lowered ox. state) and ammonia is the reducing agent (loses e-s, inc. ox. state

of N).

The reaction between ammonium and nitrate(III) (nitr i te ) ions

NH4+

(aq) + NO2-(aq) ==> 2H2O(l) + N2(g)

Here its the opposite of disproportionation where two species of an element in different oxidation states react to produce one species of

a single intermediate oxidation state.

Ox. state changes: Nitrogen in a (-3) and a (+3) state both end up in the (0) state. Oxygen at (-2) and hydrogen (+1) remain unchanged in oxidation state.

The nitrite ion acts as the oxidising agent and gets reduced (N +3 to 0, 3e's gained, decrease of 3 ox. state units)

and the ammonium ion acts as the reducing agent and gets oxidised (N -3 to 0, 3 e's lost, inc. ox. state 3 units).

The nitrite ion acts as the oxidising agent (gains/accepts e-s, lowered ox. state of N) and the ammonium ion acts as the reducing agent

(loses/donates e-s, inc. ox. state of N).

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